Chemical Bonding — Free Chemistry Review Games.

Ionic bonds form when one atom transfers electrons to another, creating oppositely charged ions that attract each other.

Covalent bonds form when two atoms share one or more pairs of electrons.

Carbon is in Group 14 and has 4 valence electrons, allowing it to form 4 covalent bonds.

The octet rule states that atoms tend to form bonds to achieve 8 electrons in their outer shell, like noble gases.

Ionic bonds typically form between metals (which lose electrons) and nonmetals (which gain electrons).

Lewis dot structures represent an atom's valence electrons as dots around its chemical symbol to visualize bonding.

A large electronegativity difference (>1.7) between atoms favors ionic bonding, while a small difference favors covalent bonding.

In a polar covalent bond, electrons are shared unequally because one atom has higher electronegativity, creating partial charges.

Metallic bonds involve a 'sea' of delocalized electrons shared among metal atoms, giving metals conductivity and malleability.

A double bond involves the sharing of two pairs (four total) of electrons between two atoms.

NaCl forms a crystal lattice with strong ionic bonds between many alternating Na+ and Cl- ions, requiring significant energy to break.

VSEPR (Valence Shell Electron Pair Repulsion) theory predicts molecular geometry by minimizing repulsion between electron pairs.

Four bonding pairs with no lone pairs arrange tetrahedrally (109.5 degree bond angles), as seen in methane (CH4).

Water's bent shape (due to two lone pairs on oxygen) and the electronegativity difference between O and H create a net dipole moment.

Intermolecular forces (like hydrogen bonds, dipole-dipole, London dispersion) act between molecules and are much weaker than covalent or ionic bonds within molecules.

A cation forms when an atom loses electrons, leaving more protons than electrons and a net positive charge. The distractor "An atom that has gained one or more electrons, giving it a negative charge" describes an anion, not a cation. Recognizing that metals typically lose electrons to form cations is essential for predicting ionic compound formulas.

An anion carries a negative charge because it has gained one or more electrons, giving it more electrons than protons. The option "A positively charged ion formed by losing electrons" actually defines a cation, which is the opposite process. Nonmetals typically gain electrons to complete their octet, becoming anions.

Oxygen is in Group 16 (VIA) of the periodic table, and elements in this group have 6 valence electrons in their outer shell. The choice "8" would represent a full octet, not oxygen's actual count before bonding. Knowing an element's group number lets you quickly determine its valence electron count for drawing Lewis structures.

A nonpolar covalent bond occurs when two atoms have very similar or identical electronegativities, so they share bonding electrons nearly equally. The distractor "Electrons are transferred completely from one atom to another" describes ionic bonding, a fundamentally different bonding type. Bonds between identical atoms, like \(\text{Cl}_2\), are classic examples of nonpolar covalent bonds.

In Lewis structures, a single line drawn between two atoms represents one shared pair of electrons, forming a single covalent bond. The option "Two shared pairs of electrons (a double bond)" is incorrect because a double bond is shown with two lines, not one. Correctly counting lines versus dots is essential for accurately interpreting bonding and lone pairs in Lewis diagrams.

Dots placed around an atom's symbol in a Lewis structure represent valence electrons that are not involved in bonding, known as lone pairs. The distractor "The number of protons in the nucleus" confuses electron representation with a completely different atomic property. Correctly distinguishing bonding electrons (lines) from lone pairs (dots) is fundamental to drawing accurate Lewis structures.

Sodium is a metal with low electronegativity that readily loses an electron, while chlorine is a nonmetal with high electronegativity that readily gains one, producing a large electronegativity difference favorable for ionic bonding. The pairing "Carbon and hydrogen" involves two nonmetals with a small electronegativity difference, which instead forms a covalent bond. As a general rule, metal-nonmetal combinations tend to form ionic bonds due to their large electronegativity gap.

Electronegativity specifically measures how strongly an atom attracts shared electrons within a covalent bond. "Ionization energy" instead measures the energy required to remove an electron from an isolated atom, which is a related but distinct concept. Electronegativity trends across the periodic table are the key tool for predicting bond polarity.

Electronegativity increases from left to right across a period because nuclear charge increases while shielding stays roughly constant, and it increases up a group because valence electrons are closer to the nucleus with less shielding. The option describing an increase "right to left across a period" reverses the correct periodic trend. Fluorine, located in the upper right of the periodic table, is the most electronegative element, which anchors this trend in memory.

Ionic compounds form rigid crystal lattices held together by strong electrostatic forces, resulting in high melting points and brittleness when force is applied and disrupts the ion alignment. The claim of "Good electrical conductivity in solid form" is false because ions are fixed in place in the solid lattice and cannot move to carry charge until melted or dissolved. This solid-versus-molten conductivity distinction is a hallmark property used to identify ionic compounds experimentally.

A triple bond consists of three pairs of electrons, or six electrons total, shared between the same two atoms, as seen in \(\text{N}_2\). The distractor "One shared pair of electrons between three atoms" incorrectly describes electron sharing among three separate atoms rather than a stronger bond between two. Triple bonds are shorter and stronger than single or double bonds because more electron density holds the nuclei together.

Magnesium is in Group 2 and loses its two valence electrons to achieve a stable noble gas electron configuration, resulting in a \(2+\) charge. The option "\(2-\)" would require magnesium to gain two electrons, which is inconsistent with its low electronegativity and metallic character. Group number on the periodic table is a reliable predictor of the typical ionic charge for main-group metals.

An electronegativity difference above roughly \(1.7\) indicates that one atom pulls the bonding electrons so strongly that essentially complete electron transfer occurs, classifying the bond as ionic. A "Nonpolar covalent bond" instead requires a very small electronegativity difference, typically less than about \(0.4\). This numeric threshold is a useful rule of thumb for classifying bond type based on electronegativity difference, though it is a continuum rather than a sharp cutoff.

The electronegativity difference between H and Cl is about \(0.9\), which falls in the intermediate range that produces unequal electron sharing and partial charges, defining a polar covalent bond. The label "Ionic" requires a much larger difference, typically greater than \(1.7\), which this pair does not reach. Calculating electronegativity differences is a practical method for classifying real bonds rather than relying on general element type alone.

Carbon dioxide requires double bonds between carbon and each oxygen so that carbon achieves an octet with four bonding pairs total and each oxygen also completes its octet with two bonding pairs and two lone pairs. "Single bonds" would leave carbon with only two bonding pairs, giving it just four electrons and violating the octet rule. Drawing correct Lewis structures often requires testing multiple bond orders to satisfy the octet rule for every atom.

Magnesium forms a \(2+\) ion by losing two valence electrons, and since each chlorine atom can only accept one electron to form a \(1-\) ion, two chloride ions are needed to balance the overall charge to neutral. The option stating chlorine "gains two electrons" is incorrect because chlorine only needs one electron to complete its octet. Balancing total positive and negative charge to zero is the key principle for determining ionic compound formulas.

In \(\text{Cl}_2\), both atoms are identical chlorine atoms with the exact same electronegativity, so the bonding electrons are shared perfectly equally, making the bond nonpolar covalent. The option "\(\text{NaCl}\)" instead has a very large electronegativity difference between sodium and chlorine, making it ionic rather than covalent. Bonds between two atoms of the same element are always nonpolar covalent because there is zero electronegativity difference.

Nitrogen contributes 5 valence electrons and each of the three hydrogen atoms contributes 1, giving a total of \(5 + 3(1) = 8\) valence electrons to distribute in the structure. The value "5" only accounts for nitrogen's electrons and ignores the hydrogens' contributions entirely. Summing valence electrons from all atoms in the formula is the essential first step before drawing any Lewis structure.

Noble gases already possess a complete octet (or duet for helium) of valence electrons, so they have no thermodynamic driving force to gain, lose, or share additional electrons through bonding. The claim that they have "very high electronegativity values that repel other atoms" is inaccurate because noble gases typically do not have meaningfully assigned electronegativity values due to their lack of bonding tendency. The octet rule explains both why most atoms bond to achieve noble gas configurations and why noble gases themselves remain largely unreactive.

Bond polarity depends on the difference in electronegativity between two bonded atoms, and since both atoms in \(\text{F}_2\) are fluorine, that difference is exactly zero, making the shared electrons distributed equally. The statement that fluorine "always creates polar bonds regardless of the partner atom" is false because polarity is a relative comparison, not an absolute property of one element alone. This example shows that even highly electronegative elements form nonpolar bonds when bonded to an identical atom.

A double bond involves two shared electron pairs instead of one, increasing the electron density between the nuclei, which pulls the atoms closer together and requires more energy to break, making it both shorter and stronger. The option claiming a double bond is "weaker and longer" reverses the correct trend seen experimentally in bond energy and bond length data. As bond order increases from single to double to triple, bond length decreases while bond strength increases.

Calcium is a metal with low electronegativity that easily loses its two valence electrons, while oxygen is a nonmetal with high electronegativity that readily accepts electrons, creating a large electronegativity difference characteristic of ionic bonding. The statement that oxygen "loses electrons to calcium, which is a nonmetal" is wrong on two counts, since oxygen gains rather than loses electrons and calcium is actually a metal. Identifying metal-nonmetal pairings with large electronegativity gaps is the standard method for predicting ionic bond formation.

The octet rule states atoms tend to bond in ways that give them eight valence electrons, so the number of bonds an atom forms typically equals the number of electrons it needs to gain or share to complete that octet. "The atomic mass of the element" has no direct bearing on bonding behavior, since bonding depends on electron configuration, not mass. This principle explains why carbon, needing four more electrons, typically forms four bonds.

As electronegativity increases from carbon to nitrogen to oxygen, the electronegativity difference with hydrogen also increases, making the bonds progressively more polar in the order \(\text{C-H} < \text{N-H} < \text{O-H}\). The option placing \(\text{O-H}\) as least polar contradicts oxygen's high electronegativity, which creates the largest difference with hydrogen among these three bonds. Comparing electronegativity values along a row of the periodic table allows students to rank bond polarities without memorizing exact numbers.

Since \(\text{NO}_3^-\) carries an overall \(1-\) charge, one extra electron beyond the neutral atoms' valence electrons must be included in the total electron count and reflected somewhere in the structure through formal charge assignment. The claim that formal charges "only apply to neutral molecules, not ions" is incorrect, since formal charge calculations are especially important for accurately representing charged polyatomic ions. Tracking total charge, including extra or missing electrons, ensures the Lewis structure correctly represents the species being drawn.

Potassium's valence electron occupies a higher principal energy level farther from the nucleus with more electron shielding, so it experiences a weaker attractive pull and is removed more easily than sodium's valence electron, making potassium more reactive in forming cations. The claim that potassium "has a higher electronegativity than sodium" is false, since electronegativity actually decreases going down a group. Reactivity trends among alkali metals in ionic bond formation are explained by decreasing ionization energy down a group.

Elements in period 3 and beyond, like sulfur, have accessible d-orbitals that allow them to expand beyond the traditional octet and accommodate more than eight electrons, as seen with the six bonding pairs in \(\text{SF}_6\). The statement that "all elements can violate the octet rule without any conditions" is inaccurate, since this expanded octet exception is generally restricted to period 3 and higher elements, not elements like carbon or nitrogen in period 2. Recognizing expanded octets as a legitimate exception, tied to orbital availability, prevents students from forcing incorrect octet-only structures for larger central atoms.

To satisfy the octet rule, sulfur in \(\text{SO}_2\) needs one double bond and one single bond along with a lone pair, distributing electrons so both sulfur and each oxygen end up with eight electrons around them through resonance-averaged bonding. The claim that sulfur "can have fewer than eight electrons" as a stable, preferred structure is generally disfavored since incomplete octets are less stable than resonance structures that satisfy the octet rule for all atoms when possible. Testing initial single-bond-only structures against total electron count and octet completion often reveals the need for double or triple bonds.

Each C=O bond is individually polar due to oxygen's higher electronegativity, but because \(\text{CO}_2\) is linear with the two identical bond dipoles pointing in exactly opposite directions, the vector sum of the dipoles cancels to zero, making the overall molecule nonpolar. The claim that carbon and oxygen "have identical electronegativities" is false since oxygen is considerably more electronegative than carbon, which is exactly why each individual bond is polar. Molecular polarity depends on both bond polarity and molecular geometry together, not on bond polarity alone.

According to Coulomb's law, electrostatic attraction between ions increases with the product of their charges, so the \(2+\) and \(2-\) charges in \(\text{MgO}\) create a much stronger lattice energy and higher melting point than the singly charged ions in \(\text{NaCl}\). The claim that "\(\text{MgO}\) has covalent bonds while \(\text{NaCl}\) has ionic bonds" is incorrect since both compounds are ionic, formed between metals and nonmetals with large electronegativity differences. Comparing ionic charge magnitude is one of the most reliable ways to predict relative lattice energy and melting point trends among ionic compounds.

Nitrogen's small atomic radius and high electronegativity allow \(\text{NH}_3\) molecules to form strong hydrogen bonds between molecules, a much stronger intermolecular force than the weaker dipole-dipole and dispersion forces present in \(\text{PH}_3\), resulting in a higher boiling point for ammonia. The claim that "phosphorus is more electronegative than nitrogen" is false, since nitrogen is actually more electronegative than phosphorus on the periodic table. Recognizing hydrogen bonding as a uniquely strong intermolecular force involving N-H, O-H, or F-H bonds is essential for predicting relative boiling points.

Aluminum's small ionic radius combined with its \(3+\) charge creates a high charge density that strongly polarizes the electron cloud of the neighboring chloride ions, distorting the electron distribution enough to introduce significant covalent character into the bond, a trend described by Fajans' rules. The statement that "all metal-nonmetal compounds are purely ionic with no exceptions" oversimplifies bonding, since real bonds exist on a continuum and highly charged, small metal cations like \(\text{Al}^{3+}\) often show partial covalent behavior. This illustrates that the ionic-covalent classification is a spectrum rather than a strict binary, especially for cations with high charge density.

Experimental data shows all three carbon-oxygen bonds in \(\text{CO}_3^{2-}\) have equal, intermediate bond lengths, which is best explained by treating the actual structure as a resonance hybrid, an average of the contributing Lewis structures rather than any single one being literally correct. The claim that the molecule "rapidly switches between forms, changing shape every instant" misrepresents resonance, since the real structure is a single, static blended hybrid rather than an oscillation between distinct structures. Resonance structures are a modeling tool to represent delocalized electron density that no single classical Lewis structure can fully capture.

Fluorine's very high electronegativity combined with its small atomic size creates a highly polarized H-F bond, concentrating a strong partial negative charge on a small fluorine atom that allows exceptionally strong hydrogen bonding between HF molecules. The claim that fluorine "forms ionic bonds with hydrogen, unlike chlorine" is incorrect, since both H-F and H-Cl bonds are polar covalent, differing only in the degree of polarity. Among hydrogen halides, only \(\text{HF}\) exhibits significant hydrogen bonding because fluorine meets the small-size, high-electronegativity criteria required for this strong intermolecular force.

In graphite, each carbon forms three sigma bonds within a planar layer, leaving one electron delocalized in a pi system that can move freely across the layer, giving graphite electrical conductivity, while diamond's four localized sigma bonds per carbon leave no free electrons for conduction. The claim that graphite "is an ionic compound while diamond is purely covalent" is false, since both allotropes are entirely composed of covalently bonded carbon atoms, just arranged differently. Comparing carbon allotropes demonstrates how bonding geometry and electron delocalization, not just bond type, determine macroscopic properties like conductivity.

Nitrogen achieves a stable octet by forming a triple bond, sharing three electron pairs between the two atoms, which pulls the nuclei closer together than the double bond in \(\text{O}_2\), which only shares two electron pairs. The claim that "oxygen atoms are smaller than nitrogen atoms" is not the primary driver here and is actually inaccurate, since atomic radius alone does not explain this trend as strongly as bond order does. Higher bond order consistently correlates with shorter, stronger bonds when comparing bonds of the same atom pair type.

Beryllium is an unusually small metal cation with high charge density that strongly polarizes the electron cloud of chloride ions, pulling electron density back toward itself and giving the bond substantial covalent character rather than being purely ionic. The blanket assumption that "all metal-nonmetal bonds are always purely ionic" ignores well-documented exceptions like beryllium and aluminum compounds, where small, highly charged cations distort anions significantly. Bond character exists on a continuum influenced by both electronegativity difference and ionic size or charge, not determined by metal-nonmetal classification alone.

Measured bond lengths in ozone show both O-O bonds are identical and intermediate between typical single and double bond lengths, which cannot be captured by any single Lewis structure with one localized double bond, requiring instead an average of two resonance structures. The option stating ozone "only has single bonds between all oxygen atoms" fails to satisfy the octet rule for the central oxygen atom, which needs one double bond contribution to complete its octet. Resonance is invoked precisely when experimental structural data conflicts with what any single Lewis structure predicts.

Lattice energy scales with the product of ionic charges according to Coulomb's law, so the doubly charged \(\text{Ca}^{2+}\) and \(\text{O}^{2-}\) ions in \(\text{CaO}\) create substantially stronger electrostatic attraction and higher lattice energy than the singly charged \(\text{K}^+\) and \(\text{F}^-\) ions in \(\text{KF}\). The claim that "potassium and calcium have identical charges" is factually wrong, since potassium forms a \(1+\) ion while calcium forms a \(2+\) ion, a key difference driving the lattice energy comparison. When comparing lattice energies, ionic charge typically has a greater impact than ionic size differences of similar magnitude.

Chlorine is located in Group 17 (VIIA) of the periodic table, meaning it has 7 valence electrons in its outer energy level. The option "8" would represent a complete octet, which chlorine reaches only after gaining one more electron through bonding. Group number provides a fast way to determine the valence electron count needed to predict bonding behavior.

When two nonmetal atoms come together, both have relatively high electronegativity and neither easily gives up electrons, so they share electron pairs to complete their octets, forming a covalent bond. "An ionic bond" instead requires one atom to transfer electrons completely to another, which typically occurs between a metal and a nonmetal. Nonmetal-nonmetal combinations are a reliable indicator that a covalent bond, rather than an ionic one, will form.

Fluorine has the highest electronegativity of any element, roughly \(3.98\) on the Pauling scale, due to its small atomic radius and strong effective nuclear charge on valence electrons. "Oxygen" is highly electronegative as well but ranks below fluorine on the standard electronegativity scale. Fluorine's position in the top-right corner of the periodic table (excluding noble gases) makes it the benchmark for maximum electronegativity.

A lone pair consists of two valence electrons that belong to a single atom and are not shared with another atom in a bond, often shown as dots in Lewis structures. The choice "A pair of electrons shared between two bonded atoms" instead describes a bonding pair, which forms a covalent bond rather than remaining unshared. Distinguishing lone pairs from bonding pairs is critical for correctly predicting molecular geometry using VSEPR theory.

Potassium bromide (\(\text{KBr}\)) forms between a metal, potassium, and a nonmetal, bromine, resulting in a large electronegativity difference and the transfer of electrons that characterizes an ionic bond. "\(\text{CH}_4\)" is composed entirely of nonmetals, carbon and hydrogen, which share electrons covalently rather than transferring them. Recognizing metal-nonmetal combinations as the classic signature of ionic bonding helps quickly classify unfamiliar compounds.

The central atom in a Lewis structure is generally the one with the lowest electronegativity, excluding hydrogen, because less electronegative atoms are typically better at forming multiple bonds to surrounding atoms. The option "Hydrogen, since it usually forms the most bonds" is incorrect because hydrogen can only form one bond, since it needs just two electrons total to satisfy its duet rule. Placing hydrogen only as a terminal (outer) atom and choosing the least electronegative other atom as central is a standard strategy for building Lewis structures.